The Chemistry of Life: Chapter 2 Study Notes

The Chemistry of Life

Introduction

The study of the chemistry of life explores the fundamental chemical principles that underlie biological processes. Understanding the elements, atoms, and molecules that compose living organisms is essential for grasping how life functions at the molecular level.

Elements Essential to Life

Major and Minor Elements in the Human Body

Living organisms are composed of a limited number of elements, most of which are found in significant proportions in the human body. These elements are crucial for biological structure and function.

• Major Elements: Carbon (C), Oxygen (O), Nitrogen (N), Hydrogen (H), Phosphorus (P), Sulfur (S), Potassium (K), Sodium (Na), Calcium (Ca), Magnesium (Mg), Chlorine (Cl), Iron (Fe)

• Trace Elements: Zinc (Zn), Manganese (Mn), Copper (Cu), Iodine (I), Selenium (Se), and others

Element	Approximate % in Body
Carbon	50%
Oxygen	20%
Nitrogen	14%
Hydrogen	8%
Phosphorus	3%
Sulfur	1%
Potassium	1%
Sodium	1%
Calcium	0.5%
Magnesium	0.5%
Chlorine	0.5%
Iron	0.2%
Trace metals	0.3%

Example: The human body contains about $4.00 worth of these elements in total.

Elements and Atoms

Definitions and Properties

All matter is composed of elements and atoms. Elements are pure substances that cannot be broken down by ordinary chemical reactions. Atoms are the smallest units of an element that retain its chemical properties.

• Element: A substance consisting of only one type of atom, with unique chemical properties.

• Atom: The smallest unit of an element, composed of subatomic particles.

• Subatomic Particles: Protons (positive charge), Neutrons (neutral), Electrons (negative charge)

Example: Carbon, sodium, and phosphorus are elements found in the body, each with distinct properties and appearances.

Atomic Structure

Subatomic Particles and Their Roles

Atoms consist of a nucleus containing protons and neutrons, surrounded by electrons in orbitals. The arrangement and number of these particles determine the atom's identity and chemical behavior.

• Proton: Positively charged particle found in the nucleus; determines atomic number.

• Neutron: Neutral particle found in the nucleus; contributes to atomic mass.

• Electron: Negatively charged particle found in orbitals around the nucleus; involved in chemical reactions.

Particle	Charge	Location
Proton	+1	Nucleus
Neutron	0	Nucleus
Electron	-1	Orbitals

Additional info: The number of protons defines the element, while the number of neutrons can vary, resulting in isotopes.

Electron Orbitals and Shells

Organization and Energy Levels

Electrons occupy specific energy levels called shells, which are made up of orbitals. The arrangement of electrons in these shells determines how atoms interact and bond.

• Shells: Energy levels around the nucleus; each shell can hold a specific number of electrons.

• Orbitals: Regions within shells where electrons are likely to be found; s, p, d, and f types.

• Electron Configuration: Electrons fill the lowest energy orbitals first (Aufbau principle).

Shell	Orbitals	Max Electrons
1st (K)	1 s-orbital	2
2nd (L)	1 s, 3 p-orbitals	8
3rd (M)	1 s, 3 p, 5 d-orbitals	18
4th (N)	1 s, 3 p, 5 d, 7 f-orbitals	32

Example: The electron configuration of lithium is $1s^2 2s^1$.

Chemical Bonding

Types of Chemical Bonds

Atoms bond to achieve stable electron configurations, usually by filling their valence shells. The main types of chemical bonds are covalent, ionic, and hydrogen bonds.

• Covalent Bond: Atoms share pairs of electrons; can be nonpolar (equal sharing) or polar (unequal sharing).

• Ionic Bond: Electrons are transferred from one atom to another, creating charged ions (cations and anions).

• Hydrogen Bond: Weak attraction between a hydrogen atom in one molecule and an electronegative atom in another.

• Van der Waals Interactions: Weak attractions due to transient electron fluctuations.

Example: Table salt (NaCl) is formed by ionic bonding between sodium and chlorine.

Electronegativity and Bond Polarity

Electronegativity Trends and Effects

Electronegativity is an atom's tendency to attract electrons in a bond. Differences in electronegativity lead to bond polarity and affect molecular interactions.

• High Electronegativity: Atoms like oxygen and nitrogen strongly attract electrons.

• Low Electronegativity: Atoms like sodium and potassium weakly attract electrons.

• Polar Covalent Bonds: Unequal sharing of electrons, resulting in partial charges.

• Nonpolar Covalent Bonds: Equal sharing of electrons.

Example: Water ($H_2O$) has polar covalent bonds due to oxygen's high electronegativity.

Properties of Water

Unique Characteristics and Biological Importance

Water's molecular structure gives it unique properties essential for life, including cohesion, adhesion, high specific heat, and its role as a universal solvent.

• Cohesion: Water molecules stick to each other via hydrogen bonds.

• Adhesion: Water molecules stick to other surfaces.

• High Specific Heat: Water absorbs and retains heat, stabilizing temperatures.

• Solvent Properties: Water dissolves polar and ionic substances easily.

Example: Water's ability to dissolve salts and sugars is crucial for cellular processes.

Acids, Bases, and pH

Water Dissociation and pH Scale

Water can dissociate into hydronium ($H_3O^+$) and hydroxide ($OH^-$) ions, affecting the pH of solutions. The pH scale measures the concentration of hydrogen ions.

• Acid: Substance that increases $H^+$ concentration; pH < 7.

• Base: Substance that decreases $H^+$ concentration; pH > 7.

• Neutral: Pure water has a pH of 7.

• Buffer: Mixture of weak acid and its conjugate base that stabilizes pH.

Equation:

$H_2O \rightleftharpoons H^+ + OH^-$

Example: Blood contains buffers to maintain a near-neutral pH.

Carbon: The Universal Adaptor

Carbon's Role in Organic Molecules

Carbon's unique atomic structure allows it to form diverse and stable bonds, making it the backbone of all organic molecules.

• Organic Compounds: Molecules containing carbon, such as lipids, carbohydrates, nucleic acids, and proteins.

• Carbon Skeleton: The chain or ring of carbon atoms forming the framework of organic molecules.

• Autotrophs: Acquire carbon from atmospheric $CO_2$.

• Heterotrophs: Acquire carbon from other organisms.

Example: Glucose ($C_6H_{12}O_6$) is an organic compound essential for energy.

Hydrocarbons and Isomers

Structure and Biological Relevance

Hydrocarbons are organic compounds containing only hydrogen and carbon. Isomers are molecules with the same chemical formula but different structures, affecting their properties.

• Alkanes: Single bonds, tetrahedral geometry.

• Alkenes: Double bonds, planar geometry.

• Alkynes: Triple bonds, linear geometry.

• Aromatic Rings: Unique alternating double-bond structure, found in nucleic acids and other biomolecules.

• Isomers: Molecules with identical formulas but different structures (e.g., cis- and trans-fats).

Example: Benzene is an aromatic hydrocarbon found in many biological molecules.

Functional Groups in Organic Molecules

Types and Biological Roles

Functional groups are specific groups of atoms attached to carbon skeletons that confer unique chemical properties and reactivity to organic molecules.

• Amino Group (-NH2): Found in amino acids.

• Carboxyl Group (-COOH): Found in organic acids.

• Phosphate Group (-PO4): Found in nucleic acids and ATP.

• Sulfhydryl Group (-SH): Found in cysteine and methionine; important for protein folding.

Example: The phosphate group is essential for the structure of DNA and RNA.

Hemoglobin and Gas Exchange

Role of pH and Functional Groups

Hemoglobin is a protein in red blood cells that carries oxygen. Its affinity for oxygen and carbon dioxide is influenced by pH and the functional groups of its amino acids.

• Oxyhemoglobin: Form with high affinity for oxygen.

• Deoxyhemoglobin: Form with reduced affinity for oxygen, higher affinity for carbon dioxide.

• Bohr Effect: Acidic conditions (lower pH) promote release of oxygen and uptake of carbon dioxide.

Equation:

$HCO_3^- + H^+ \rightleftharpoons H_2CO_3 \rightleftharpoons CO_2 + H_2O$

Example: Increased acidity in tissues causes hemoglobin to release oxygen where it is needed most.

Regulation of Breathing

Control by the Nervous System

Breathing rate is regulated by nerve centers in the brain that sense carbon dioxide levels in the blood. When $CO_2$ levels rise, the body is triggered to inhale and exhale to maintain homeostasis.

• Central Chemoreceptors: Detect changes in $CO_2$ and pH in the blood.

• Inhalation Trigger: Increased $CO_2$ leads to increased acidity, stimulating breathing.

• Homeostasis: Maintained unconsciously by the brain.

Example: Holding your breath increases $CO_2$ and acidity, causing the urge to breathe.